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chem notes week 11

by: Andrea Scota

chem notes week 11 CHE 106 - M001

Andrea Scota
GPA 3.7
General Chemistry Lecture I
R. Doyle

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Chem notes for week 11
General Chemistry Lecture I
R. Doyle
Class Notes
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This 0 page Class Notes was uploaded by Andrea Scota on Wednesday November 11, 2015. The Class Notes belongs to CHE 106 - M001 at Syracuse University taught by R. Doyle in Fall 2015. Since its upload, it has received 38 views. For similar materials see General Chemistry Lecture I in Chemistry at Syracuse University.

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Date Created: 11/11/15
Pink mentioned in class Chem Notes Week 1 1 TEXTBOOK CHAPTER 8 sections 84 88 Basic Concepts of Chemical Bonding contd Bond Polarity and Electronegativity 84 0 Bond polarity measure of how equallyunequally the electrons in any covalent bond are shared 0 electrons are shared equally 0 one atom exerts a greater attraction for bonding electrons that the other 0 Uses electronegativity to estimate whether a given bond is polar covalent or nonpolar covalent or ionic F 0 Greater electronegativity the greater its ability to attract electrons to itself 0 An atom with very negative electron affinity and a high ionization energy both attracts electrons from other atoms and resists having them taken away meaning they are highly electronegative o Fluorine is the greatest Electronegativity values o The difference in electronegativities of bonded atoms can be used to determine the polarity of a bond the grater the electronegative difference the more polar the bond 0 Polar molecule one whose centers of positive and negative charge do not coincide A polar molecule has a positive side and a negative side 0 Polar interactions account for may properties of liquids solids and solutions Whenever 2 electrical charges of equal magnitude but opposite sign are separated by distance a is established o Magnitude of a dipole is measured by denoted with u o Q is the equalopposite charges r is distance o Dipole moment increases as the magnitude of Q increases and as r increases Larger the dipole moment the more polar a bond 0 Measures in diybes D 334 X 10quot30 Columbmeters 0 Q 160 X 10quot19 C electrostatic charge 0 r distance in angstroms Any diatomic molecule X Y in which X and Y have different electronegativities is polar Most bonding interactions lie between the extremes of covalent and ionic bonding While true that binding between metalnonmetal is ionic there are some exceptions when differences in electronegativity of atoms is relatively small or when oxidation state of metal becomes large Drawing Lewis Structures 85 0 Lewis structures help us understand bonding in many compounds 0 1 Sum valence elctrons from all atoms taking into account overall charge anion 1 electron to total for each negative charge cation 1 electron from total for each positive charge 0 2 Write the symbols for the atoms show which atoms are attached to which and connect them with a single bond a line represents 2 electrons The central atom is usually the least electronegative o 3 Complete the octets around all the atoms bonded to central atom o 4 Place any left over electrons on the central atom o 5 o of any atom in a molecule is the charge the atom would have if each bonding electron pair in a molecule were shared equally between 2 atoms 0 In generally the dominant Lewis structure will have low formal charges with any negative formal charged residing on the more electronegative atoms 0 The formal charges do not represent the real charges Resonance Structures 86 0 Sometimes a single dominant Lewis structure is inadequate to represent a single molecule or ion 0 Both or all Lewis structures are they are called The molecule Important in describing bonding in ozone and organic molecule benzene Exceptions to the Octet Rule 87 Three main types of exceptions o 1 Moleculespolyatomic ions containing an odd of electrons o 2 Molecules and polyatomic ions in which an atom has o 3 Molecules and polyatomic ions in which an atom has Lewis structures with more than an octet of electrons are observed for atoms in the third row and beyond in the periodic table Strengths and Lengths of Covalent Bonds 88 The stability of a molecule is related to the strengths of its covalent 0 Related to change in enthalpy enthalpy change AH for the breaking of a particular bond in one mole of a gaseous substance Average bond enthalpies can be determined for a wide variety of covalent bonds Strength of covalent bonds increases with the of electron pairs shared between 2 atoms Bond enthalpy is energy is always required to breach chemical bonds therefore energy is always released when a bond forms Greater bond enthalpy stronger bond consistent with the bond being stronger as the number of bonds increasing


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